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Chemistry for Biology Students (CHEM0010) Notes - Structure, Bonding and Self-Assembly of Biological Structures CA$11.89   Add to cart
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Chemistry for Biology Students (CHEM0010) Notes - Structure, Bonding and Self-Assembly of Biological Structures
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  • Course
  • Chemistry for Biology Students (CHEM0010)
  • Institution
  • University College London (UCL)

Explore foundational chemistry with these specialized notes tailored for Year 1 students in the Chemistry for Biology Students (CHEM0010) module at University College London. Dive into the intricacies of structure, bonding, and self-assembly of biological structures, deciphering key concepts like a...

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Document information

  • November 30, 2023
  • 12
  • 2020/2021
  • Class notes
  • Professor andrea sella
  • All classes

Subjects

  • atom
  • electrons
  • orbitals
  • covalent bonding
  • lewis
  • vsepr
  • valence bonding
  • molecular orbitals
  • van der waals
  • hydrogen bonds
  • self assembly

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  • University College London (UCL)
  • Unknown
  • Chemistry for Biology Students (CHEM0010)
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A1: Structure of the Atom, Electrons and Orbitals
A1: Structure of the Atom, Electrons and Orbitals
 Atomic structure
o Subatomic particles
 Electron – mass = 0, charge = -1, spin = 0.5
 Proton – mass = 1, charge = +1, spin = 0.5
 Neutron – mass = 1, charge = 0, spin = 0.5
o Nucleus
 Proton + neutron
 Atomic number (Z) = number of protons in the nucleus
 Relative atomic mass (Mr) = number of protons + neutrons
o Element
 Fixed number of protons
 Isotope = varying number of neutrons
 Mass spectrometer – measures mass of isotopes
o Shows peaks for mass of each isotope + relative abundance
 Electron
o Electron configuration – determines chemical reactivity
 Occupy orbitals = have fixed energy values
o Rules:
 Aufbau’s ‘building up’ principle – lowest energy orbitals are occupied first
 Pauli principle – only 2 electrons occupy each orbital
o Electrons have – wave-particle duality
 Wave-particle duality – a quantum object is both a wave and a particle
 Electrons can be described by a wave function:
o Wave function – mathematical distributions of electrons in terms of position
and time
 Heisenberg Uncertainty Principle
 Position and momentum cannot simultaneously be determined
o The energy of an electron is known (due to spectroscopy) > therefore the
position is unknown
o Orbital – region where an electron is most probably located
 Quantum numbers – describe orbital shapes and size (electrons have 4 quantum numbers)
 Principle quantum number (n) –1 st quantum number (shell)
o Shell (n)
 Subshell – within each shell
o How big the orbital is / how far an electron occupying it is from the nucleus
 Angular quantum number (I) – 2nd quantum number (shape of orbital)
o I = 0, … n-1 (n = principle quantum number)
o Tells you the shape of the orbital (s, p, d, f)
 S: I = 0
 Each s subshell – has 1 s orbital
 P: I = 1
 Each p subshell – has 3 p orbitals
 D: I = 2
 Each d subshell – has 5 d orbitals
 Magnetic quantum numbers (m1) – 3rd quantum number (orientation of orbital)
o m1 – -I, … +I
 E.g. if I = 1, m1 = -1, 0, +1
o Tells you the orientation of the orbital
 x = -1

, A1: Structure of the Atom, Electrons and Orbitals
 y=0
 z = +1
 Spin quantum number (ms) – 4th quantum number (spin)
o Electrons have a spin – property of electron is related to angular momentum
 ms = -0.5, +0.5
 ms – +0.5 = clockwise rotation
 ms – -0.5 = anticlockwise rotation
 Pauli exclusion principle – no 2 electrons in an atom can have the same 4 quantum numbers
 Max 2 electrons can share each orbital
o Same value of n
o Same value of I
o Same value of m1
o Different (ms) = different spin
 Summary of quantum numbers

n (shell) I (shape) Orbital notation m1 (orientation) Orbital names
1 0 1s 0 1s
2 0 2s 0 2s
1 2p -1, 0, +1 2px, 2py, 2pz
3 0 3s 0 3s
1 3p -1, 0, +1 3px, 3py, 3pz
2 3d -2, -1, 0, +1, +2 3dxy, 3dyz, 3dzx,
3dx2-y2, 3dz2
 Shape of orbitals (I)
o Wave functions:
 Order of energies of orbitals – orbital further from
the nucleus = higher energy
 Node = 0% chance of finding an electron
o More nodes = higher energy
 Radial distribution function – how the probability of
finding an electron varies with distance from the
nucleus
 1s2 electrons shield 2s1 electron from the
nuclear charge
o But: 2s radial distribution function
has a hump closer to the nucleus
than 1s
 Resulting in an energy boost
for 2s1 = can penetrate than
1s2 orbital + spends more
time closer to the nucleus
 2s1 closer to nucleus
= experiences a
stronger attraction –
energy is lowered
 2p6 electrons radial distribution function = low probability of finding electron close
to nucleus
o 2p6 electrons do not penetrate 1s2 orbital as much = better shielded from
nuclear charge
 Higher energy
 Hund’s rule of maximum multiplicity

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